Arsenate and antimonate adsorption competition on 6-line ferrihydrite monitored by infrared spectroscopy

Applied Geochemistry 61 (2015) 224–232

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Arsenate and antimonate adsorption competition on 6-line ferrihydrite monitored by infrared spectroscopy Tim Muller a,1, Dave Craw b, A. James McQuillan a,⇑ a b

Department of Chemistry, University of Otago, P.O. Box 56, Dunedin 9054, New Zealand Department of Geology, University of Otago, P.O. Box 56, Dunedin 9054, New Zealand

a r t i c l e

i n f o

Article history: Available online 18 June 2015 Editorial handling by M. Kersten

a b s t r a c t Arsenate and antimonate are water-soluble toxic mining waste species which often occur together and can be sequestered with varying success by a hydrous ferric oxide known as ferrihydrite. The competitive adsorption of arsenate and antimonate to thin films of 6-line ferrihydrite has been investigated using primarily adsorption/desorption kinetics monitored by in situ attenuated total reflectance infrared (ATR-IR) spectroscopy on flowed solutions containing 103 and 105 mol L1 of both species at pH 3, 5, and 7. ICP-MS analysis of arsenate and antimonate adsorbed to 6-line ferrihydrite from 103 mol L1 mixtures in batch adsorption experiments at pH 3 and 7 was carried out to calibrate the relative surface concentrations giving rise to the IR spectral absorptions. The kinetic data from 103 and 105 mol L1 mixtures showed that at pH 3 antimonate achieved a greater surface concentration than arsenate after 60 min adsorption on 6-line ferrihydrite. However, at pH 7, the adsorbed arsenate surface concentration remained relatively high while that of adsorbed antimonate was much reduced compared with pH 3 conditions. Both species desorbed slowly into pH 3 solution while at pH 7 most adsorbed arsenate showed little desorption and adsorbed antimonate concentration was too low to register its desorption behaviour. The nature of arsenate which is almost irreversibly adsorbed to 6-line ferrihydrite remains to be clarified. Ó 2015 Elsevier Ltd. All rights reserved.

1. Introduction Arsenic and antimony exist widely in the environment, mainly from anthropogenic activities. Mining brings arsenic minerals such as arsenopyrite (FeAsS) to the surface where oxidation forms arsenite (As(III)) and eventually the (As(V)) species such as arsenate (AsO3 4 ) which is very effectively sequestered by a hydrous ferric oxide known as ferrihydrite (Roddick-Lanzilotta et al., 2002; Goldberg and Johnston, 2001; Carabante et al., 2009; Loring et al., 2009). Antimony in low oxidation states may also be found in mines in minerals such as stibnite (Sb2S3) which undergoes surface oxidation on exposure to atmospheric oxygen to release antimonite (Sb(III)) and form Sb(V) species such as antimonate (Sb(OH) 6 ). The adsorption of arsenate to ferrihydrite across quite a wide pH range is an important component of waste water remediation at many mining sites (Mohan and Pittman, 2007). Antimonate is also adsorbed to ferrihydrite and this is more pronounced under acidic conditions with markedly less adsorption under neutral conditions (McComb et al., ⇑ Corresponding author. Tel.: +64 3 4797924. E-mail addresses: [email protected] (T. Muller), jmcquillan@chemistry. otago.ac.nz (A.J. McQuillan). 1 Present address: Babbage Consultants Ltd, New Zealand. http://dx.doi.org/10.1016/j.apgeochem.2015.06.005 0883-2927/Ó 2015 Elsevier Ltd. All rights reserved.

2007; Ilavsky, 2008; Vithanage et al., 2013; Essington et al., 2013; Guo et al., 2014). In mine water situations where both arsenate and antimonate are present in comparable concentrations they compete for adsorption sites at the ferrihydrite surface (Milham and Craw, 2009; Wilson et al., 2010; Kolbe et al., 2011). The adsorption affinities (from adsorption isotherms) of arsenate and of antimonate on ferrihydrite have been obtained under certain conditions as single adsorbate species (Roddick-Lanzilotta et al., 2002; McComb et al., 2007; Essington et al., 2013). Such data may be useful to predict the outcome of mixed adsorption situations if all adsorption sites are equivalent, a condition which is difficult to determine. However, direct evaluation of the relative propensities of arsenate and antimonate to adsorb to ferrihydrite from a solution containing both adsorbates relates better to practical contexts such as mine waste water treatment. Water and rock analyses at the Reefton gold mine, New Zealand have established that treatment of water containing comparable concentrations of As and Sb by ferrihydrite adsorption removes most of the arsenic but little of the antimony (Milham and Craw, 2009) (see Fig. 1). Ritchie et al. (2013) have shown that sorption and co-precipitation with iron (hydr)oxides are important attenuation pathways in Alaskan watersheds and that elevated antimony levels extend further downstream from lodes than elevated arsenic levels.

T. Muller et al. / Applied Geochemistry 61 (2015) 224–232

Fig. 1. Covariation of dissolved As and Sb at Reefton gold mine, New Zealand. Open squares are waters in the processing plant in which arsenopyrite and stibnite are concentrated, and these waters, like much of the ore, have approximately equal As and Sb (1:1 line indicated). Processing waters are passed through a treatment plant which generates hydrous ferric oxide (HFO)/ferrihydrite from ferric chloride.

ATR-IR spectroscopy, with internal reflection prisms coated with thin particle films of the solid substrate, has become established as a valuable in situ means of establishing the molecular nature of adsorbate–adsorbent interactions and measuring adsorbed species concentrations with variations in solution pH and adsorbate concentration (Biber and Stumm, 1994; McQuillan, 2001; Lefevre, 2004; Mudunkotuwa, 2014). The method was first used to investigate the in situ adsorption of mine waste arsenate to hydrous ferric oxide by Roddick-Lanzilotta et al. in 2002 and it was concluded that arsenate forms a bidentate inner-sphere complex which is strongly bound in neutral to acid solutions (Roddick-Lanzilotta et al., 2002). A subsequent such study by Carabante et al. (2009) with ferrihydrite in D2O drew attention to the impact of a slow carbonate desorption rate in natural systems on the kinetics of arsenate adsorption. Their use of D2O provided improved signal quality of arsenate absorptions at low wavenumber. Loring et al. (2009) have questioned the bidentate nature of arsenate adsorption at iron oxide surfaces and argued that monodentate coordination at goethite surfaces is more consistent with XRD, EXAFS and IR data. Several other studies of arsenate adsorption at iron oxide surfaces (Fuller et al., 1993; Sabur et al., 2015) have concluded that a mixture of monodentate and bidentate adsorption occurs. McComb et al. (2007) studied the adsorption of antimonate to hydrous ferric oxide using ATR-IR. They found that antimonate is mainly adsorbed by an outer-sphere interaction under neutral conditions, making it vulnerable to desorption by anion exchange. However, in acidic conditions an inner-sphere coordinative adsorption develops to render it more resistant to desorption. Adsorption affinities can be derived from adsorption isotherms measured using ATR-IR spectroscopy, by allowing an adsorbent/adsorbate system enough time to reach equilibrium in a series of increasing concentrations, but the process is slow and the precision in any obtained equilibrium constant is generally low. Adsorption affinities can in principle be obtained more rapidly via adsorption and desorption kinetics data. However, a recent ATR-IR analysis of oxalate adsorption on TiO2 by this approach (Young and McQuillan, 2009) showed that reaction enthalpy effects may need to be addressed before reliable adsorption rate constants and hence equilibrium constants can be obtained. In that study (Young and McQuillan, 2009) the data for the slower desorption process, which is less thermally disturbing than the adsorption, showed clear temporal distinction between first order decays of inner- and outer-sphere adsorbed oxalate signals. Thus adsorption/desorption kinetics experiments can provide

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semiquantitative comparisons of adsorbate systems involving more than one type of adsorption and on a shorter timescale than required for equilibrium experiments. These experiments (Young and McQuillan, 2009) suggested that the same approach could be conveniently used to analyse adsorption from a mixture of adsorbates. In the present work the adsorption/desorption kinetics approach in combination with in situ ATR-IR spectroscopy using single reflection prisms has been employed to evaluate the competitive adsorption of arsenate and antimonate to 6-line ferrihydrite at different pH and adsorbate concentrations. Additionally, batch adsorption experiments under conditions corresponding to those used in the adsorption/desorption kinetics experiments were carried out to calibrate the absorbance of characteristic arsenate and antimonate IR peaks and assess the relative amounts adsorbed. It is shown that arsenate adsorbs strongly to 6-line ferrihydrite across a wider pH range than antimonate which presents more cause for environmental concern in mine waste treatment away from acidic conditions. 2. Materials and methods 2.1. Materials Potassium hexahydroxyantimonate, KSb(OH)6 (Merck, Pro Analysi); sodium arsenate heptahydrate, Na2HAsO47H2O (Univar, 98%) and iron (III) chloride hexahydrate, FeCl36H2O (Unilab, 98%) were used as received. All solutions were prepared using 18 MX cm resistivity deionised water (Millipore, Milli-Q). Adjustments to pH for ATR-IR experimental solutions were made using dilute solutions prepared from NaOH (Merck, 99%) and HCl (AnalaR, 12.0 ± 0.5 mol L1) and no pH buffers were used to prevent any interference with adsorption. For quantitative adsorption experiments, HNO3 (AnalaR, 15.8 ± 0.1 mol L1) was used instead of HCl to minimise chloride ion related ICP-MS interference. The ferrihydrite for IR spectroscopy was prepared from FeCl36H2O according to the method of McComb et al. (2007). 1 mL of 0.7 mol L1 FeCl3 solution was added drop-wise to 50 mL of boiling water and stirred vigorously for 5 min. The resulting dark red suspension was allowed to cool before its acidity was reduced by dialysis over 24 h. The ferrihydrite was stored for up to about one week in small HDPE bottles. A transmission XRD of the freeze-dried material was recorded between polymer films on a Stoe (Darmstadt) diffractometer using a molybdenum source and a Mythen 1 K (Dectris) 1 strip detector. This XRD (see Supplementary Data Fig. S1) confirmed that the iron oxide was 6-line ferrihydrite as previously determined by Carabante et al. (2009). An iron oxide thin film for each ATR-IR spectroscopy experiment was prepared by depositing 200 lL of the aqueous suspension on a ZnSe prism, covering 1 cm2, which was then dried for 60 min using a water pump vacuum (40 kPa). Washing each film with aqueous NaOH (3  103 mol L1) for 10 min prior to adsorption experiments removed most adsorbed carbonate. All solutions flowed over the films were presaturated with N2. 2.2. Infrared spectroscopy of solutions and adsorbed species A Digilab FTS-4000 infrared spectrometer with Digilab Merlin™ software (version 3.4) was used to collect all infrared spectra from 64 co-added scans at a resolution of 4 cm1. For the adsorption experiments the spectrometer was fitted with a Harrick FastIR™ ATR accessory having a single 45° reflection ZnSe prism and fitted via an O-ring with a Teflon flow cell having an internal volume of

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0.22 mL. A flow rate of 2.0 mL min1 was used for all experiments. As the spectrometer was purged with dry air containing varying amounts of carbon dioxide, the spectra exhibit CO2 peaks at 2350 cm1 of varying intensity. Spectra of solution species were obtained by flowing 0.05 mol L1 solutions of arsenate or antimonate over the ZnSe prism. Due to its insolubility in acidic conditions a solution spectrum of antimonate at this concentration at pH 3 could not be obtained. Spectra of adsorbed species were obtained by flowing solutions of arsenate and/or antimonate at the appropriate pH and concentrations over an iron oxide film deposited on the prism. The highest concentration used for mixed adsorbate solutions was 103 mol L1 in both sodium arsenate and potassium antimonate. This concentration is below the reported solubility for sodium antimonate (Blandamer et al., 1974) of 3.3  103 mol L1 at 25 °C and no precipitation or cloudiness was observed in experimental solutions. Reference spectra for solution spectra or for spectra of adsorbed species were from water flowing over the prism or the iron oxide-coated prism, respectively, at the same pH as the sample solution. Spectra from each adsorption experiment were collected every 2–10 min for 60 min of the process. Subsequently, the desorption process was examined by flowing water across the iron oxide coated prism with pH adjusted to that used in the adsorption process. Desorption was followed for at least 120 min, with spectra taken at least every 10 min. Kinetics of adsorption and desorption were analysed by plotting the absorbance of arsenate- and antimonate-related peaks at 825 and 1110 cm1, respectively, versus time. In replicate experiments relative absorbance of arsenate and antimonate peaks were consistent within about 5% while there was 10–20% variation in spectral absorbance between experiments due to variability in particle film deposition.

the initial adjustment. Changes in volume due to pH adjustments were obtained from mass measurements and found to be less than 0.4 mL in all cases. Volume changes were accounted for in calculations where necessary. The solutions were stirred for 24 h using magnetic stirrer bars. After this, the hydrous ferric oxide was removed from suspension by centrifugation (Heraeus Sepatech Labofuge 200; 5300 rpm; 1 h), and the supernatant was decanted into sample vials. The small amount of excess water remaining in the iron oxide after centrifugation was considered insignificant. Nitric acid (5 mL of 0.32 mol L1) and hydrochloric acid (3 mL of 1.0 mol L1) were then added to the centrifuge tube and the resulting suspensions were magnetically stirred until the iron oxide dissolved completely (5 h). These solutions were then decanted into sample vials. For ICP-MS analysis both the supernatant and dissolved film solutions were then serially diluted 1:840 (1:40, then 1:21) with 0.32 mol L1 quartz-distiled HNO3 into 5 mL polypropylene vials. Each sample was analysed for 57Fe, 75As and 121Sb using an Agilent 7500ce Inductively Coupled Plasma Mass Spectrometer (ICP-MS). American Public Health Association (APHA) Method 2005, 3125B was used for all samples. An internal standard containing 45Sc, 103Rh, 115In, 159Tb and 209Bi was used in all samples and blanks. 57Fe, 75As and 121Sb counts were taken as a ratio to 45 Sc, 103Rh and 115In, respectively, ensuring consistency across the full mass range. Helium was used as a collision gas (flow rate = 4.3 mL min1) to eliminate interferences. Of most concern were Ar–Cl background interferences which can lead to false arsenic counts. Using this technique, these were reduced to 250 s1 for a 1% HCl solution, less than 1% of the lowest sample count (31,595 s1). All uncertainties stated for experimental results are the standard deviation of replicate values.

2.3. Batch adsorption experiments

3. Results and discussion

To obtain quantitative adsorption data to accompany the ATR-IR results, batch adsorption experiments were performed. The ferrihydrite was prepared very similarly to that prepared for IR spectroscopy. Dropwise addition of 1.43 mL of 0.7 mol L1 FeCl3 to 75 mL of boiling water precipitated the hydrous ferric oxide which was then vigorously stirred for 10 min. The resulting suspension was dialysed for 24 h and the volume was then made up to 100 mL in a volumetric flask. This 102 mol L1 Fe(III) suspension was separated into two 50 mL aliquots, which were adjusted to pH 3 and pH 7. In 15 mL plastic centrifuge tubes, 9 mL of the iron oxide suspensions were mixed with 1 mL of a solution containing sodium arsenate and potassium antimonate, both at 1.17  102 mol L1, giving arsenate and antimonate concentrations of (1.17 ± 0.07)  103 mol L1, a concentration close to that used in the ATR-IR adsorption/desorption experiments with 103 mol L1 solutions. To prevent precipitation of sodium antimonate the 1.17  102 mol L1 mixture was prepared while warm (40 °C). ICP-MS results for arsenate and antimonate recovery (94 ± 2% and 101 ± 4%, respectively) confirmed that no significant precipitation had taken place. Three replicate samples were studied at each pH. Additionally, two experimental blanks (at pH 3 and pH 7) were prepared, in which 1 mL of deionised water was used in place of the arsenate/antimonate solution. After addition of the arsenate + antimonate solution (or deionised water in the case of blanks), fine pH adjustments were made as necessary. As no buffers were used, pH measurements of each sample were made throughout the experiment, and pH adjusted where necessary. No experimental sample was found to be more than 0.4 from the desired pH after

3.1. IR spectra of arsenate and antimonate in aqueous solution It is useful to obtain the solution spectra of adsorbate species because the nature of adsorbed species may often be recognised by comparison with the spectra of solution species which are generally better characterised. Solution ATR-IR spectra for arsenate (0.05 mol L1 Na2HAsO4) at pH 3 and 7 are shown in Fig. 2. H3AsO4 has pKa (25 °C) values (Greenwood et al., 1984) of 2.2, 6.9 and 11.5. Thus H2AsO 4 is expected to be dominant over 2 H3AsO4 at pH 3.0, with H2AsO 4 and HAsO4 being almost equally significant at pH of 7.0. At pH 3, arsenate exhibits prominent spectral peaks at 908 and 877 cm1. These observations agree well with those of Myneni et al. (1998) and earlier reports which assigned these peaks to the As–O antisymmetric and symmetric stretching bands, respectively, of the C2v symmetry molecule. Smaller peaks observed here at 735 and 748 cm1 appear to correspond to the poorly-defined features observed respectively at 730 cm1 (pD 4), 750 cm1 (pH 3), and 759–766 cm1 (pH 5.7) by other authors (Carabante et al., 2009; Myneni et al., 1998). Myneni et al. (1998) assigned these bands predominantly to the As–OH symmetric and antisymmetric stretching modes of H2AsO-4. The broad weak absorption at 1200 cm1 has been attributed to As–O–H deformation vibrations (de Portilla, 1976). At pH of 7, the most prominent arsenate spectral absorption is a fairly broad peak at 858 cm1 with a shoulder at 903 cm1 and a weak absorption at 1194 cm1. The 858 cm1 band has been attributed to the absorptions of overlapping antisymmetric and symmetric As–O stretch modes of HAsO2 at 865 and 846 cm1, respectively (Myneni 4 et al., 1998). The 903 cm1 shoulder and 1194 cm1 absorption show that some H2AsO 4 is present at this pH, as expected.

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Fig. 2. ATR-IR spectra for 0.05 mol L1 solutions of: (A) arsenate at pH 3 (offset by absorbance of 0.010); (B) arsenate at pH 7 (offset by absorbance of 0.005); and (C) antimonate at pH 7 (not offset). The background for each spectrum was from water adjusted to the same pH.

An antimonate solution at pH of 7 is almost entirely Sb(OH) 6, based on the pKa for antimonic acid HSb(OH)6 of 2.72 (Baes and Mesmer, 1976). The IR spectrum of 0.05 mol L1 KSb(OH)6 solution at pH 7 is dominated by a broad peak at 1016 cm1 arising from the in-plane O–H deformation mode of Sb(OH) 6 which has an octahedral SbO6 core (McComb et al., 2007). At pH 3 due to the presence of some HSb(OH)6 there was insufficient solubility to enable a solution IR spectrum to be obtained. Further Sb(OH) 6 absorptions (McComb et al., 2007) not shown here, are found at 3126 cm1 in the O–H stretching region and an out-of-plane O–H deformation at 767 cm1.

3.2. IR spectra of species adsorbed to 6-line ferrihydrite from 103 mol L1 arsenate and antimonate solution In this study adsorption competition was considered at pH of 3, 5, and 7 for solutions containing both arsenate and antimonate at 103 mol L1 and at 105 mol L1 where the latter concentration corresponds more closely to typical mine effluent conditions. Experiments were carried out with an initial adsorption phase of 1 h followed by a desorption phase of several hours. Spectra resulting at the end of the adsorption phase are first reported followed by the corresponding kinetics data. Fig. 3 shows ATR-IR spectra taken 1 h after adsorption to iron oxide from pH 3, 5 and 7 solutions containing 103 mol L1 of both arsenate and antimonate. Adsorbed arsenate is evident in a peak which varies between 837 and 818 cm1 in the pH range between 3 and 7. This variation in peak wavenumber with pH has been

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Fig. 3. ATR-IR spectra of adsorbed species obtained from flowing solutions of pH 3, 5 and 7 containing arsenate and antimonate both at 103 mol L1 over ferrihydrite films for 1 h. The background for each spectrum was obtained from water at the same pH flowing over the film. The pH 5 and pH 3 spectra are offset by 0.05 and 0.10, respectively, on the absorbance scale.

observed previously (Roddick-Lanzilotta et al., 2002; Carabante et al., 2009) and suggests a change in degree of protonation of the adsorbed arsenate. An additional or alternative explanation for this peak shift is an increase in the proportion of monodentate to bidentate adsorbed arsenate species with increased pH. This explanation is considered in more detail later. Adsorbed antimonate is detected in the broad OH stretching band in the 3100 cm1 region, and the corresponding OH bending bands at 1100 cm1, and at 760 cm1 (see Supplementary Data Fig. S3). Only minor differences in adsorbed arsenate and antimonate spectral features were observed from the mixed species solutions compared with spectra from single species adsorption. See the Supplementary Data spectra in Figs. S2 and S3 from adsorbed single species at pH 3 and 7 for 103 mol L1 solutions under identical conditions to those employed for the mixed adsorbate systems. These single adsorbate spectra concur with those previously published by this group (Roddick-Lanzilotta et al., 2002; McComb et al., 2007). However, some shifts in the 1100 cm1 antimonate and 830 cm1 arsenate peaks were observed which may arise from differences in interactions between like and unlike adsorbed species and differences in adsorption sites. Kolbe et al. (2011) have shown that the presence of antimonate has no influence on the adsorption of arsenate to akaganeite at pH 7 but the adsorption of antimonate decreases in the presence of arsenate. The use of single internal reflection prisms allows interesting spectral features in the OH stretching region to be included which are not so accessible using multiple reflection accessories. Noticeable spectral features arising from the 1 h of adsorption

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include the pronounced absorption loss consisting of negative peaks at 3593 and 3400 cm1. The absorption loss at 3400 cm1 corresponds to that of bulk water being displaced by the adsorbed species and is accompanied by the less prominent water bending mode loss 1640 cm1. A similar absorption loss to that of the distinctive peak at 3593 cm1 has previously been attributed to loss of interfacial OH species with antimonate adsorption to iron oxide3 and was accompanied by a corresponding shift to lower wavenumber in D2O. A possible alternative explanation may be the loss of ‘‘free OH’’ bonds of water at the iron oxide–water interface when adsorbed species displace interfacial water but there is only a small concentration in the bulk and its absorption is centred about 3650 cm1 (Max and Chapados, 2010). The bulk water absorption loss concomitant with adsorption is greatest at pH 3 and this influences the apparent wavenumber of the antimonate peak which drops to 3045 cm1 compared with 3115 cm1 at pH 7. The absorption loss features in the 1600–1300 cm1 region arise from displacement of adsorbed carbonate due to the adsorption of arsenate and antimonate. Without stringent measures to avoid contact with the atmosphere, evidence for carbonate and bicarbonate adsorption is commonly observed in metal oxide IR spectra (Dobson and McQuillan, 1997; Hausner et al., 2009) and needs no further discussion here. Additionally, such displacement of adsorbed carbonate will invariably be part of arsenate and antimonate adsorption to HFO in practical contexts. For arsenate, As–O stretching bands were observed in Fig. 3 at 837 cm1 for pH 3, 825 cm1 for pH 5, and 818 cm1 for pH 7 and the absorbance of these bands after 60 min was 0.0306, 0.0296 and 0.0231, respectively. For antimonate, with bands at 1109 cm1 for pH 3, 1106 cm1 at pH 5 and little significant intensity at pH 7, the corresponding band absorbance after 60 min was 0.0105 and 0.0050. This data appears to indicate stronger arsenate adsorption at lower pH as is commonly found for anion adsorption on metal oxides. However, the reproducibility between replicate experiments of about 20% for adsorbed species absorbance needs to be considered. Also, this conclusion is not so simply reached when the speciation of the adsorbed arsenate species changes with pH, as evident in the increase in peak wavenumber at lower pH. Nevertheless, the retention of considerable adsorbed arsenate under neutral conditions is the outstanding feature of this data. The shape of the antimonate peak near 3100 cm1 is influenced by the adjacent water absorption changes resulting in an apparent shift in the absorption maximum to lower wavenumber with lower pH. Water absorption is very strong in this region and even small changes in interfacial water structure have relatively large spectral impact which is not yet well understood. The 1100 cm1 antimonate peak which is less influenced by adjacent absorptions of other species has been used as an indicator of relative amount of adsorbed antimonate, although it must be assumed that any minor arsenate absorption from the wing of the weak As–O–H deformation band at 1200 cm1 is insignificant . Thus the absorbance of the arsenate at 830 cm1 and of antimonate 1100 cm1 have been used to monitor the relative amounts of these individual species adsorbed. It has also been assumed that the 760 cm1

antimonate band makes no contribution to the arsenate peak absorbance 830 cm1. Assuming the Beer Law applies in this context, and this is generally true (Mudunkotuwa, 2014), relating measured absorbance to surface concentrations requires molar absorption coefficients (e) of the respective bands. These molar absorption coefficients are presently unknown but their ratios can be determined through batch adsorption experiments. 3.3. Relative amounts of arsenate and antimonate adsorbed to 6-line ferrihydrite from batch adsorption experiments We have used batch adsorption experiments at pH 3 and 7 and ICP-MS analysis to obtain the relative amounts of arsenate to antimonate adsorbed to 6-line ferrihydrite under comparable conditions to those used for the IR adsorption/desorption experiments in which both adsorbates were at 103 mol L1. It has been assumed that the relative concentrations of adsorbed arsenate and antimonate would remain the same in the different measurements. The results of the ICP-MS analysis of the samples from the batch adsorption experiments are shown in Table 1. The batch adsorption data shows a surprisingly large amount of adsorbed arsenate and antimonate relative to the amount of iron in each batch. Such a high adsorbed-species-to -adsorbent amount ratio have been previously explained by the high specific surface areas and very small particles known for ferrihydrite suspensions (Fuller et al., 1993; Waychunas et al., 1993). The possibility of ferric arsenate or scorodite precipitation in such adsorption experiments has been previously raised and discounted (Waychunas et al., 1993) and the IR spectra from the present work have given little support for the presence of these solids which have strongest absorptions at 740 and 790 cm1, respectively (Gomez et al., 2009; Swedlund et al., 2015). For the pH 3 batch adsorptions the arsenate and antimonate solution concentrations remaining in solution of 0.66 and 0.49  103 mol L1, respectively, were marginally less than the 103 mol L1 of both adsorbates used for the spectral data reported in Fig. 3. Nevertheless, the batch adsorption data shows that adsorbed antimonate coverage is 1.6 times that of arsenate under these conditions, which dispels contrary suggestions about relative amounts adsorbed based simply on relative peak absorbance at 1109 and 837 cm1. Table 1 data along with the measured absorbance of each of these bands, assuming Beer law validity, enables the following ratio of molar absorption coefficients to be derived.

e837 ðAsÞ  4:7 e1109 ðSbÞ where em~ (i) is the molar absorption coefficient for adsorbed species i ~ cm1 in the infrared spectrum. at m Thus the molar absorption coefficient for the antisymmetric As– O stretch mode of the adsorbed arsenate species at 830 cm1 is much larger than that of the adsorbed antimonate OH in-plane deformation mode at 1100 cm1. In the batch experiments an adsorption time of 24 h was allowed in comparison with 1 h adsorption in the IR spectroscopy experiments and the adsorption

Table 1 ICP-MS analysis of batch experiments for adsorption of arsenate and antimonate ion to ferrihydrite. pH

cAs/104 mol L1

nAs/106 mol

cSb/104 mol L1

nSb/106 mol

nSb/nAs

3 7

6.6 ± 0.4 8.3 ± 0.5

4.2 ± 0.1 2.9 ± 0.1

4.9 ± 0.2 11.26 ± 0.04

6.7 ± 0.4 0.99 ± 0.01

1.6 ± 0.1 0.34 ± 0.02

Table ci are equilibrium concentrations remaining in solution and ni are equilibrium amounts adsorbed. In each batch, amount of iron present nFe was 8.8  105 mol and initial solution concentration of both species was 1.17  103 mol L1. All errors are total SD. Arsenate and antimonate recoveries averaged 98 ± 5% with the experimental blank typically <1% of sample concentration.

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behaviour beyond 1 h would need to be considered in further such comparisons. The results of the batch adsorption experiments at pH 7 show considerably less adsorbed antimonate relative to adsorbed arsenate, in contrast to the results at pH 3. This difference is a reflection of the strong pH dependence of antimonate adsorption but weak pH dependence of arsenate adsorption in the pH range studied. In these batch experiments the arsenate and antimonate solution concentrations of 0.83 and 1.13  103 mol L1 are closer to the 103 mol L1 of the spectra experiments. The spectral data appear to indicate that very little antimonate is adsorbed at pH 7 and does not provide a very good basis for precise analysis of the relative amounts of antimonate and arsenate adsorbed. However, the batch adsorption data at pH 7 indicates the amount of adsorbed antimonate is about a third of that for adsorbed arsenate. The antimonate spectral band is quite small and its magnitude may be influenced by adjacent carbonate loss bands seen in this spectrum and account for this discrepancy.

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while the corresponding 105 mol L1 mixture at pH 3 has a 831 cm1 peak absorbance of 0.0121. The lower amount adsorbed is expected from the isotherm data for arsenate (Roddick-Lanzilotta et al., 2002) and antimonate (McComb et al., 2007) single species adsorption to ferrihydrite. The drop in arsenate surface concentration is much less than the 100 fold drop in solution concentration due to surface saturation being approached in the 105–104 mol L1concentration range. For the 105 mol L1 experiments the amount of adsorbed arsenate appears to be relatively constant across the pH range taking the film reproducibility into account. Overall, the amount of adsorbed antimonate relative to the amount of adsorbed arsenate is remarkably similar to that exhibited in the 103 mol L1 spectral data. This appears to imply very similar adsorption isotherms and adsorption affinities for these different adsorbates. Furthermore for pH 3 the relative amount of adsorbed antimonate is again greater than that of adsorbed arsenate, in this case by about 45%. 3.5. Adsorption/desorption kinetics for 103 mol L1 arsenate and antimonate on 6-line ferrihydrite

3.4. IR spectra of species adsorbed to 6-line ferrihydrite from 105 mol L1 arsenate and antimonate solution The adsorption behaviour at 105 mol L1 is more pertinent to practical mine effluent conditions. The ATR-IR spectra arising from 1 h of adsorption to ferrihydrite from solutions containing 105 mol L1 in both arsenate and antimonate at pH of 3, 5, and 7 are shown in Fig. 4. It is noticeable from the absorbance scale that the amount of adsorbed arsenate and adsorbed antimonate is much lower at this concentration compared to the 103 mol L1 data in Fig. 4. For example, adsorbed arsenate for the 103 mol L1 mixture at pH 3 has a 837 cm1 peak absorbance of 0.0306

Fig. 4. ATR-IR spectra of adsorbed species obtained from flowing solutions of pH 3, 5 and 7 containing arsenate and antimonate both at 105 mol L1 over ferrihydrite films for 1 h. The background for each spectrum was obtained from water at the same pH flowing over the film. The pH 5 and pH 3 spectra are offset by 0.025 and 0.050, respectively, on the absorbance scale.

The in situ ATR-IR studies carried out with flowed solutions have considerable potential for study not only of adsorption kinetics but also of the corresponding desorption kinetics. Equilibrium constants for adsorption such as those derived for Langmuir adsorption (KL) are in favourable contexts determinable from the ratio of adsorption rate constant to desorption rate constant (ka/kd) (Young and McQuillan, 2009). Thus differences in adsorption/desorption kinetics can provide insights about adsorption affinities. However, when two adsorbates compete for adsorption sites, as often occurs in practical situations, analysis of the resultant data is more complex. Fig. 5 shows the variation of adsorbed arsenate peak absorbance at 825 cm1 and adsorbed antimonate absorbance at 1100 cm1 during the 60 min adsorption process from 103 mol L1 solution followed by the desorption process into solutions at the corresponding pH but devoid of adsorbate species. For comparison purposes the measured peak absorbance at 60 min of arsenate adsorption and of antimonate adsorption have been normalised. The wavenumber for measured absorbance at each pH is given in Fig. 5 caption. For the 103 mol L1 data at pH of 3 in the absorbance-normalised Fig. 4 comparison there is rapid adsorption of both arsenate and antimonate with very similar time dependence to almost plateau after 60 min, at which time the solution flow was switched to that without the adsorbates and the desorption started. At pH 3 the arsenate is predominantly H2AsO 4 in solution with some H3AsO4 while the antimonate is less predominantly Sb(OH) 6 with a more significant proportion of undissociated HSb(OH)6. Thus the ferrihydrite positive surface charge might be expected to play a greater role in assisting the adsorption rate of arsenate than that of antimonate. The relative amounts of adsorbed antimonate to adsorbed arsenate being 1.6 after 60 min may suggest that other factors such as the relative population of specific adsorption sites for inner-sphere antimonate and arsenate are influential as adsorption rate depends on site population (area) (Young and McQuillan, 2009). In comparison to the adsorption data, the desorption data shows a significant differentiation between the desorption curves with antimonate at first appearing to be more rapidly desorbing. Considering that desorption rate is proportional to surface concentration (coverage) (Young and McQuillan, 2009), and that initial coverage of antimonate is a little higher than that of arsenate under these conditions, it appears that there are fairly similar desorption rate constants for the two species and therefore little differentiation in adsorption affinities. Both desorption plots have more rapid although different rates of absorbance loss in the first

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Fig. 5. Temporal variation of normalised IR absorbance for arsenate peak (D) and antimonate peak () for pH 3 at 837 and 1109 cm1, respectively, pH 5 at 825 and 1106 cm1, respectively, and for pH 7 at 818 cm1 during 1 h adsorption to ferrihydrite from flowed solution containing arsenate and antimonate both at 103 mol L1 followed by desorption. The kinetic data came from the same experiments giving the Fig. 3 spectra. Data from pH 5 and pH 3 spectra are offset on absorbance scale. Absorptions associated with antimonate were not observed significantly at pH 7.

hour with subsequently more gradual decays at similar rates in the second hour of desorption. This observation will be revisited later in this section. At pH of 5, where surface charge is less and H2AsO 4 and Sb(OH) 6 dominate in solution, the antimonate adsorption is noticeably slower than the corresponding arsenate data to reach up to a plateau. Using the molar absorption coefficient ratio deduced for pH 3 conditions and applying it also to the pH 5 spectra with the slightly shifted peaks at 1106 and 825 cm1 gives an estimated surface coverage of antimonate relative to that of arsenate of 0.74. Thus there must be other factors such as site preference influencing this difference. The corresponding desorption rates are quite different initially with the apparent arsenate desorption rate being about one third that of antimonate for the first 30 min. Considering both species are adsorbing to and desorbing from comparable surface concentrations there must be significantly different factors influencing the adsorption behaviour of antimonate compared with that of arsenate. At pH of 7 there is little antimonate adsorption evident from the spectra and only the arsenate adsorption/desorption data is shown in Fig. 4. The adsorption curve is comparable to that for arsenate at lower pH and, as for the lower pH data, there is a faster desorption component accompanied by a slower one. For this pH the slower desorption is very slow and after 2 h there is no further desorption with about 76% of the adsorbed arsenate remaining from that at the start of the desorption. A closer examination of the evolution with time of the spectral data from the pH of 7 adsorption and desorption processes has

Fig. 6. IR spectra of arsenate adsorbed to ferrihydrite from solution at pH of 7 containing arsenate and antimonate both at 103 mol L1 from the same experimental data shown in Figs. 2 and 3. The red spectrum is from 1 h after commencement of adsorption, the blue spectrum is after 1 h of desorption, with both spectra having the same background from conditions prior to adsorption. The black absorption loss spectra were recorded during the first 1 h of desorption with the background from the spectrum of adsorbed arsenate after 1 h of adsorption (the red spectrum). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)

revealed details of the different adsorbed arsenate species. Fig. 6 is of the spectral region containing the strong arsenate absorption band. The spectra with positive absorbance shown in Fig. 6 include the peak of adsorbed arsenate after 1 h of adsorption as well as that from after the initial 1 h of desorption into pH 7 solution, after which a more gradual desorption rate was observed. These typically broad adsorbed species bands appear to consist of just two overlapping peaks at about 822 and 863 cm1. This would suggest the two absorption spectra are not significantly different as a result of the 1 h desorption. However, Fig. 6 absorption loss spectra showing spectral changes over the initial 1 h of desorption reveal that the species lost in this period has a somewhat different spectrum from that of the predominant and more persistently adsorbed arsenate species under these conditions. The absorption loss spectra in Fig. 6 showing peaks at 885 and 832 cm1 retains similarities to that of the adsorbed species after 1 h adsorption while having better resolved peaks which are shifted to higher wavenumber, corresponding to a more weakly bound adsorbed species with a higher As–O average bond order. Shifts to higher wavenumber of strong As–O antisymmetric stretch absorptions occur with increased protonation of solution arsenate species (Myneni et al., 1998) so the desorption at pH 7 of a monoprotonated species is a possible explanation. Alternatively the presence of a significant concentration of inner-sphere monodentate bound arsenate may account for this more readily desorbing component of the adsorbed layer while a second component of more strongly bound bidentate arsenate desorbs much more slowly. Loring et al. (2009) have argued that on goethite at pH 6.5 a monodentate inner-sphere adsorbed arsenate bound

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formation (Okada et al., 1982). However, EXAFS studies on freeze-dried samples have concluded that bidentate interactions occur between antimonate and ion oxides (Scheinost et al., 2006; Guo et al., 2014). The sampling conditions of the present in situ ATR-IR spectroscopy work differ significantly from those during the EXAFS measurements from which the sample aqueous phase had been removed. Additionally, there appear to be no reports of hexahydroxyantimonate ions acting as bidentate ligands in metal-ligand coordination complexes, which is the usual comparison with coordinative adsorption to metal oxides. At pH of 7 there was no significant adsorbed antimonate to consider as was also observed in the 103 mol L1 data of Fig. 4. The much slower adsorption rates for this 105 mol L1 data compared with that for 103 mol L1 is also evident and largely reflects the expected first order dependence of adsorption rate on adsorbate concentration. At pH of 5 the adsorption rate for antimonate is surprisingly greater than that of arsenate even when the antimonate is more readily desorbed, which is in contrast to the corresponding behaviour for the 103 mol L1 solution mixture.

4. Conclusions

Fig. 7. Temporal variation of normalised absorbance for 825 cm1 arsenate peak (D) and 1100 cm1 antimonate peak () during 1 h adsorption to ferrihydrite from flowed solution containing arsenate and antimonate both at 105 mol L1 followed by desorption. The kinetic data came from the same experiments giving the Fig. 6 spectra. Data from pH 5 and pH 3 spectra are offset on absorbance scale. Absorptions associated with antimonate were not observed significantly at pH 7.

additionally by hydrogen bonding is predominant and that band shifts with pH primarily reflect changes in the hydrogen bonding. A recent study of arsenical adsorption on hematite supported simultaneous monodentate and bidentate adsorption of arsenate (Sabur et al., 2015) at pH of 7. A similar model for arsenate adsorption on ferrihydrite at pH 7 has some support from this data. 3.6. Adsorption/desorption kinetics for 105 mol L1 arsenate and antimonate on 6-line ferrihydrite The adsorption/desorption kinetics data for the 105 mol L1 arsenate and antimonate mixture is shown in Fig. 7. The outstanding feature of this data is the very low desorption rate of adsorbed arsenate under all pH conditions with the desorption rate decreasing from pH of 3–7 where there is almost negligible desorption over 3 h. This low arsenate desorption rate at this lower solution concentration is a reflection of the low proportion of the more weakly adsorbed (possibly monodentate) arsenate present in comparison with that of the more strongly adsorbed (possibly bidentate) arsenate. In general a low adsorption affinity results in low adsorbed species coverage at concentrations sufficient for stronger adsorbing species to achieve much higher coverage. In contrast, more than 20% of the adsorbed antimonate is desorbed within 1 h at pH 3 and 5 and by 3 h about 35% is desorbed for both pH conditions. The broadness of the adsorbed antimonate IR absorptions precludes the usual spectral assignment of binding denticity but the lability of adsorbed antimonate suggests a monodentate species, which is consistent with reported Fe–O–Sb bond

A hitherto unexplored ATR-IR spectroscopic adsorption/ desorption kinetics approach to competitive adsorption has been used to analyse the adsorption of arsenate and antimonate to ferrihydrite which is of relevance to mine waste treatment. The spectral data from this analysis showed well-separated prominent peaks of adsorbed arsenate 825 cm1 and of adsorbed antimonate 1100 cm1 which have provided relative surface concentrations of these adsorbates, by utilising relative molar absorption coefficients derived from batch adsorption experiments. For adsorption from pH 3 solutions there was a greater surface concentration of adsorbed antimonate than of adsorbed arsenate for both 103 and 105 mol L1 solutions. However, for pH 7 solutions, the adsorbed arsenate surface concentration remained relatively high while that of antimonate was much reduced compared with pH 3 conditions. This difference in adsorption affinity to ferrihydrite of arsenate and antimonate under para-neutral conditions is also evident in the desorption kinetics where an inner-sphere adsorbed arsenate species was remarkably persistent under solution flow. While the structures of outer-sphere species are closely related to those of the solution species from which they arise the nature of inner-sphere adsorbed arsenate and antimonate species are still debated and there have been relatively few spectroscopic studies in quest of this knowledge. The nature of arsenate which becomes almost permanently adsorbed to ferrihydrite, particularly under neutral conditions, and plays an extremely valuable role in environmental arsenic sequestration, warrants further investigation. Acknowledgements We acknowledge funding from the University of Otago and the NZ Ministry of Business, Innovation and Employment (MBIE). We are grateful to Jacob Shephard for XRD recording at University College London. Thanks are due to David Savory and Suzie Warring for assistance with graphical material.

Appendix A. Supplementary material This data includes the 6-line ferrihydrite XRD as well as the single species adsorbed arsenate and adsorbed antimonate IR spectra at pH 3 and 5 for 103 mol L1 aqueous solutions. Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.apgeochem.2015.06.005.

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